Riassunto esame Chimica, prof. Saracco, libro consigliato Chemistry: The Molecular Nature of Matter and Change di Silberberg

How close can molecules approach each other?

When we measure the distance between two Cl nuclei in Cl2 we obtain two different values.

• Bond length and covalent radius. The shorter distance, called bond length, is between to

bonded atoms in the same molecule.

• Van der Waals distance and radius. The longer distance is between two nonbonded atoms in

adjacent molecules. Here intermolecular attractions balance electron cloud repulsions.

Ion-dipole forces:

When an ion and a nearby polar molecule attract each other, an ion-dipole force results. A perfect

example is when an ionic compound dissolves in water. The attraction between the ions and the

oppositely charged poles of the water molecules are stronger that the attractions between he ions

themselves.

Dipole-dipole forces:

The polar molecules in liquids and solids lie near each other, and their partial charges give rise to

dipole-dipole forces. These forces depend on the magnitude of the molecular dipole moment.

The Hydrogen bond:

A special type of dipole-dipole force arises between molecules that have a H atom bonded to a

small, highly electronegative atom with lone electron pairs (specifically N, O or F). in these cases

H-bonds form. The small sizes of the atoms are essential because:

• The atoms are so electronegative that their covalently bonded H is highly positive.

• The lone pair on N, O and F of another nearby molecule can come close to the H.

This kind of bond has a profound impact on many systems.

Polarizability and induced dipole forces:

A nearby electric field can induce a distortion in the electron cloud of an atoms, pulling electron

density towards a positive pole.

• For a nonpolar molecule the distortion induces a temporary dipole moment.

• For a polar molecule in enhances the dipole moment already present.

The source of the electric field can be the charge of an ion or the partial charge of a polar molecule.

How easily and electron cloud can be distorted is called its polarizability. Smaller atoms are less

polarizable than larger ones because their electrons are closer to the nucleus. Thus we observe

several trends.

• Polarizability increases down a group as size increases.

• It decreases across a period as Zeff increases making the atoms smaller and holding electron

more tightly.

• Cations are less polarizable than their parent atoms because they are smaller. The opposite is

true for anions.

Polarizability affects all molecular forces.

London (dispersion) forces:

The intermolecular force responsible for the condensed states of nonpolar substances is the

dispersion (London) force. Dispersion forces are present between all particles because they are

caused by motion of electrons in an atom.

• Source: At any instant there may be more electron on one side of an atom than the other,

which gives the atom an instantaneous dipole. When molecules are close together this dipole

induces a dipole in its neighboring molecules and they attract each other. At low

temperatures these attractions are able to keep the molecules together.

• Prevalence: They are the only forces existing between nonpolar molecules, but contribute to

the energy of attraction in all substances because they exist between all particles. The

dispersion force is the dominating intermolecular force.

• Relative strength: The relative strength of dispersion forces depends on the polarizability of

the particles, s they are weak for small particles, but stronger for larger particles.

Polarizability depends on the number of electrons, which is correlated to the molar mass. As

molar mass increases dispersion forces increase and so do boiling points.

• Effect of molecular shape: For a pair of nonpolar substances with the same molar mass, a

molecular shape that has more area over which electron clouds can be distorted allows for

stronger attractions.

The properties of the liquid state:

Liquids have regions that are orderly one instant and random the next. Despite this complexity,

many macroscopic properties such as surface tension, capillarity and viscosity are well understood.

Surface tension:

Intermolecular forces have different effects of a molecule at the surface compared to a molecule in

the interior.

• An interior molecule is attracted by others on al sides.

• A surface molecule is only attracted by other below and to the sides, so it experiences a net

attraction downward.

This means a surface molecule tends to move into the interior. For this reason a liquid surface tends

to have the fewest molecules and as a consequence the smallest area possible. The only way to

increase the area is for molecules to move up by breaking attractions in the interior, which require

energy. The surface tension is the energy required to increase the surface area and has unit J/m^2.

In general the stronger the forces are between particles the more energy it takes to increase the

surface area, so the grater the surface tension.

Capillarity:

The rising of a liquid against the pull of gravity is called capillary action or capillarity. It is a

competition between the intermolecular forces within the liquid and those between the liquid and

the tube walls (adhesive forces).

• Water in glass: Water molecules form adhesive H bonds with the O atoms in the glass. As a

result a thin film of water creeps up the wall. At the same time cohesive H bonding forces

between water molecules, which give rise to surface tension, make the surface taut.

Adhesive and cohesive forces combine to raise the water level and produce the concave

meniscus.

• Mercury in glass: The cohesive forces among the mercury atoms are metallic bonds, much

stronger than the dispersion adhesive forces between mercury and glass. As a result the

liquid pulls away from the walls. At the same time the surface atoms are being pulled toward

the interior by mercury's high surface tension, so the levels drops. These combined forces

produce the convex meniscus.

Viscosity:

Viscosity is the resistance of a fluid to flow, and it results from intermolecular attractions.

Temperature and molecular shape influence viscosity.

• Effect of temperature: Viscosity decreases with heating. Faster moving molecules overcome

intermolecular forces more easily.

• Effect of molecular shape: Small, spherical molecules make little contact and pour easily.

Liquids containing longer molecules have higher viscosity.

The uniqueness of water:

Solvent properties of water:

The great solvent power of water results from its polarity and H bonding abilities.

• It dissolves ionic compounds through ion-dipole forces.

• It dissolves polar nonionic substances by H bonding.

• It dissolves nonpolar atmospheric gases to a limited extent through dipole-induced dipole

and dispersion forces.

Thermal properties of water:

• Specific heat capacity: Because water has so many strong H bonds, its specific heat capacity

is higher that any common liquid.

• Heat of vaporization: Numerous strong H bonds give water a very high hat of vaporization.

Surface properties of water:

Hydrogen bonding is also responsible for water's high surface tension and high capillarity.

The unusual density of solid water:

The large spaces within ice make the solid form of water less dense than the liquid form. This

behavior has major effects in nature.

• Surface of ice lakes: Aquatic life would not survive from year to year.

• Nutrient turnover: Alternation of sinking and rising water with different densities distributes

nutrients and dissolved oxygen.

• Soil formation: Repeated freeze thaw stress cracks the rock and helps produce sand and soil.

The solid state :structure, properties and bonding:

Structural features of solids:

We can divide solids into two broad categories:

• Crystalline solids have well defined shapes because their particles occur in a well defined

arrangement.

• Amorphous solids have poorly defined shapes because their particles lack an orderly

arrangement.

The crystal lattice and the unit cell:

The particles in a crystal are packed tightly in an orderly three dimensional array. Consider the

particles identical spherical atoms and imagine a point at the center of each. The collection of

particles forms a regular pattern called the crystalline lattice.

The unit cell is the smallest portion that gives the crystal if repeated in all directions.

There are 7 crystal systems and 14 types of unit cells that occur in nature, but the solid states

of a majority of metallic elements, some covalent compounds and many ionic compounds occur as

cubic lattices. A key parameter is the coordination number: the number of nearest neighbors of a

particle. There a three types of cubic unit cells.

• In the simple unit cell the centers of eight identical particles define the corners of a cube.

The particles do not touch diagonally along the cube faces or through its center. The

coordination number of each particle is 6.

• In the body centered cubic unit cell identical particles lie at each corner and at the center of

the cube. Those at the corner do not touch each other, but they all touch the center one. Each

particle is surrounded by eight others, so the coordination number is 8.

• In the face centered cubic unit cell identical particles lie at each corner and at the center of

each face, but not in the center of the cube. The coordination number is 12.

How many particles make up a unit cell? For particles of the same size, the higher the coordination

number, the higher the number of particles in a given volume. A particle at a corner or face is hare

by adjacent unit cells: the particle at the corner is shared by eight unit cells, so each one has 1/8 of

the particle.

• A simple cubic unit cell contains 1 particle.

• A body centered cubic unit cell contains 2 particles.

• A face centered cubic unit cell contains 4 particles.

Packing efficiency and the creation of unit cells:

Unit cells result from the ways atoms pack together. Packing efficiency is the percentage of the

total volume occupied by the spheres themselves.

• In the simple cubic unit cell large diamond shaped spaces are formed. The spheres only

occupy 52% of the cell volume. This is a very inefficient way to pack atoms.

• The body centered cubic unit cell uses space more efficiently. Spheres are placed on the

diamond shaped spaces of the lower layer. Its packing efficiency is 68%.

• The hexagonal and face centered cubic unit cells: in these spheres are packed most

efficiently. The large diamond shaped spaces become smaller triangular shaped spaces and

the second layer is placed upon these spaces.

◦ Hexagonal unit cell: If the third layer of spheres is placed over spaces that have another

sphere under them, such that they are in he position corresponding to that of the spheres

on the first layer, we have hexagonal closest packing based on the hexagonal unit cell.

◦ Face centered unit cell: If we place the third level of spheres over sp